CHEM1910 Inorganic Lecture 7 & 8 PDF

Group 2 Elements - Alkaline Earth Metals Ø These elements show similar reactivity but are less reactive than group 1 metals Ø Show a tendency to lose 2 electrons to form divalent cations – M 2+ Ø Metallic character increases going down the group Ø They are denser than group 1 elements ns2 Ø Most ber...

Group 2 Elements - Alkaline Earth Metals Ø These elements show similar reactivity but are less reactive than group 1 metals Ø Show a tendency to lose 2 electrons to form divalent cations – M 2+ Ø Metallic character increases going down the group Ø They are denser than group 1 elements ns2 Ø Most beryllium (Be) compounds and some magnesium (Mg) compounds are covalent rather than ionic. E.g. BeCl2 and MgH2 Properties of Alkaline Earth Metals q Melting point and boiling point decreases down the group q First ionization energy decreases down the group q Atomic radius and ionic radius increases down the group Zeff , Atomic radii and Ionization Energy Ø Nuclear charge increases when going from group 1 to group 2. Hence group 2 metals are smaller and have higher 1st ionization energies than group 1 metals in the same period. Reactions of Group 2 Elements The properties of s-block metals are the result of their loosely held valence electrons q Undergo the same kinds of redox reactions as the alkali metals but they are less reactive q Since first ionization is larger than group 1, elements of group 2 are less reactive q General reactivity trend is: qChemistry dominated by the +2 oxidation state (Be is an exception) qCompounds are stabilized by the relatively low values of 1st and 2nd ionization energies and high lattice enthalpy values. Reactions of Group 2 Elements The properties of s-block metals are the result of their loosely held valence electrons Ø Alkaline Earth metals react with halogens to yield ionic halide salts, MX2: Ø They react with oxygen to form oxides, MO: Be and Mg are relatively unreactive toward oxygen at room temperature. Both burn with a brilliant white glare when ignited by a flame. Ca, Sr, and Ba are reactive enough that they are stored under oil. Sr and Ba form peroxides MO2. Reaction with Water Group 2 metals react with water to yield hydrogen gas and a metal hydroxide, M(OH)2 Ø Be forms acidic solutions Ø Mg only reacts at temperatures above 100 °C Ø Ca and Sr reacts slowly at room temperature Ø Ba reacts vigorously Reaction with Water Group 2 metals react with water to yield hydrogen gas and a metal hydroxide, M(OH)2 Beryllium is an exception. In aqueous solutions Be becomes coordinated with four water molecules, [Be(H2O)4]2+ These solutions are acidic because of hydrolysis The high charge density on the beryllium ion polarizes the O-H bond. This makes it easier to remove a hydrogen from water. Stability of Group 2 Oxoanions Decomposition temperature of group 2 carbonates, nitrates and hydroxides Diagonal relationship between Lithium and Magnesium ► Lithium and magnesium exhibit similar properties such as: 1. Both metals react with N2 to produce nitrides. 2. Both metals burn in air to form oxides and not peroxides. 3. Their carbonates and nitrates decompose when heated to produce oxides. 4. Both ions (Li+ and Mg2+) are more easily hydrated than the others in their respective groups. 5. Both form an extensive network of organometallic compounds that have a large degree of covalency in the M-C bonds. Lithium and Beryllium; normal anamolies Lithium is the only group 1 metal that reacts with N2 to give a nitride. On burning on air Li give LiO2 whereas other group 1 elements give peroxides and superoxides LiNO3 decomposes on heating to form Li2O. Other members of group 1 decompose to form nitrites The carbonate, fluoride and hydroxides are far less soluble in water for lithium than for other group 1 metals Beryllium metal does not react with water Beryllium is amphoteric – It reacts with acids and alkalis. Beryllium compounds have much greater covalent character, e.g. BeCl2 BeO does not react with water Be2+ is acidic in aqueous solutions. Both Li and Be are small cations with high charge density and this is the genesis of these anomalies CHEM1910 Introductory Chemistry III Lecture 7b P block elements P Block Chemistry – What we will discuss Ø Describe the physical properties and general aspects of the Group 13, 14 and 15 elements such as their: o periodic trends o metallic character o oxidation states o structures and bonding in the oxides and halides trends in bond enthalpies Ø Discuss terms such as catenation and allotropy in relation to some non-metallic elements. Ø Explain clearly the structural differences and bonding of certain main group non-metallic elements and their allotropes: boron, carbon, nitrogen and phosphorous. Ø The chemistry of oxygen and it’s compounds Reading – Chemistry3 Chapter 27 Introduction to P-Block Elements o P-block elements are found in groups 13 – 18 in the periodic table. o They all have the valence electron configuration ns2 npx, where x is between 1 and 6 o The only block of the periodic table that contains elements that are metals, nonmetals, and metalloids. (Recall how IE, Zeff, c, and atomic radii change as we go across periods and down groups). o Changes in the physical properties of the elements can be related to periodicity. o P-block elements form compounds with ionic lattices, covalent network structures, and discrete molecules. o Ionic bonding between p-block elements leads to compounds with high covalent character, as the difference in electronegativity (c ) is low. o VSEPR can be used to predict the structure of discrete covalent molecules, and partial charges on the atoms can be used to glean information on how a molecule is expected to react. Periodic Trends of P-Block Elements Trends in Metallic Character Metals are found to the left of each period, and non-metals to the right. Metallic character increases down a group. Recall that in metals, valence electrons are delocalized over the metallic structure and are not associated with any one metal atom. Ionization energy (IE) increases from left to right across each period with increasing Zeff. (Few exceptions due to electronic configuration). As IE increases, it becomes harder to remove valence electrons from each atom, leading to a decrease in metallic character as we go across the period from left to right. Metallic character increases down a group with decreasing ionization energy. As a group is descended, the valence electrons lie in orbitals with successively higher principal quantum numbers that are further away from the nucleus. Oxidation States Ø Most p-block elements can adopt more than one oxidation state. The oxidation states of the elements are assigned using their electronegativities. c Si = 1.90 F Si F -1 c N = 3.04 -1 +4 F -1 F -1 -1 F c F = 3.98 +3 N -1 F F -1 c H = 2.20 N +1 H -3 H +1 H +1 o The sign of the oxidation state of an element is dependent on the atom it is bonded to o Due to this, conventional oxidation states of covalently bonded molecules are referred to as formal oxidation states. Common formal oxidation states for p-block elements Ø The maximum oxidation state increases across each period from left to right. Periodic Trends in Oxidation States – Highest Oxidation States of oxides and fluorides Period 2 Period 3 B C N O F Boron (III) B2O3 BF3 carbon (IV) CO2 CF4 Nitrogen(V) / (III) N2O5 NF3 Oxygen (II) OF2 Fluorine (-I) OF2 Al Si P S Cl Aluminum (III) Al2O3 AlF3 Silicon (IV) SiO2 SiF4 Phosphorus (V) Sulfur (VI) SO3 BF6 Chlorine (VII) Cl2O7 ClF5 [Cl(V)] P4O10 PF5 Ne Ar o The maximum oxidation state increases across a period, though expansion of the octet is not possible for the second period elements o Maximum oxidation state of the oxide corresponds to involvement of all valence s and p electrons in bonding. ClF7 does not exist. Difficult placing seven fluorides around chloride. Iodine heptafluoride (IF7) is known. Oxidation States - Trends Ø For groups 13-15 the maximum oxidation state is easily obtained and is observed by all the elements in the group. Ø Going down the group the maximum oxidation state is not as important and the oxidation state two lower than the maximum dominates. E.g. B(III) is common while B(I) is not, while for Tl, the +1 oxidation state is more common. Ø Inert pair effect – ns2 valence electrons do not participate in bonding. (Pg 1207 Chemistry3) Oxides Fluorides Some allotropes of p-block elements B C O a-rhombohedral, brhombohedral, b-tetragonal boron Diamond, graphite, fullerenes, coke dioxygen, ozone P S White, red, black Many catenated rings, amorphous AS Se Yellow, metallic/ grey, black Red(a, b, g), grey, black Sn Sb Grey, White blue, yellow, black Bi amorphous, crystalline You are not required to know the allotropes in grey B12 Unit – The building block of Boron Allotropes § Icosahedral (20-faced) structure based on clusters of 12 atoms. § Each boron atom shares eight electrons despite being so small and having only three electrons to contribute. § The three-dimensional network of bonds makes boron very hard and chemically unreactive. § Incorporating boron fibers into plastics produces a composite material: § Tougher than steel. § Lighter than aluminum. § Used in aircraft, missiles, and body armor Allotropes of Carbon Carbon Nanotubes Ø Discovered in 1997 by Japanese Scientist Sumio Lijima Ø Fullerene related cylindrical carbon molecules Single Multi Wall nanotube Wall nanotube (SWNT) (MWNT) Ø Consists of single layer carbon sheets that rolled up and connected into cylinders Ø Diameter of the cylinder is typically between 5-15 nm. Ø Can be opened at both ends but usually capped portion of the fullerene molecule (ie pentagonal rings). Ø All carbons sp2 hybridized Carbon Nanotubes Ø Nanotubes align themselves into “ropes” held together by Van Der Waals forces. Ø Extraordinary strength and unique electrical properties have found them applications in nano- electronics, optics and other materials applications Nitrogen and oxygen both exists as diatomic gases Allotropes of Groups 15 and 16 Heavier elements such as sulfur and phosphorus, however, forms the normal number of single electron pair bonds as expected from their electronic structures (2 and 3 respectively). This leads to these elements existing either as discrete molecules or chain structures which are more stable than diatoms. For heavier elements in period 3 and lower p-p bonding is less effective. Allotropes Phosphorus White Phosphorus, P4 Ø Soft, waxy, toxic solid – discrete tetrahedral Ø Low melting point (44 oC) and soluble in non-polar solvent Ø Standard state of phosphorus Ø P-P distance = 221 pm and P-P-P bond angle = 60o Ø Small angles indicate that molecule is strained, i.e. total energy of the 6 P-P bonds in P4 is less than the total energy of 6 P-P bonds formed by phosphorus atoms having normal bond angles. Ø Molecule is metastable under normal conditions and is stored under water away from oxygen – most reactive form of phosphorus. It is easily ignited and has been used in bombs. Red Phosphorus Ø Heating P4 in the absence of oxygen at 300 oC for several days leads to the formation of red phosphorus Ø Red phosphorus is less reactive than white phosphorus. Ø Ring strain is released when upon heating one of the P-P in the tetrahedron is broken and connects to another to form polymeric red phosphorus. Ø Several crystalline forms exists Melting point = 579 oC Less soluble and less reactive than P4 Black Phosphorus Black P flaky solid graphite-like appearance. Thermodynamically the most stable form. Least reactive. Obtained by heating P4 at 12,000 atm or at 220-370 oC for 8 days in presence of Hg. Structure is a double layer lattice of P6 rings. P-P within a layer = 220 pm. Interlayer – 390 pm. polymorphs Heating a-sulfur at 95.5 oC gives b-sulfur (monoclinic sulfur, mp = 119 oC), and heating both a- and b-sulfur between 112 – 119 oC gives g-sulfur. All these molecules have the molecular formula S8, with the same covalent linkages. These are not allotropes but polymorphs. Allotropes Oxygen a special element of the p-block Oxygen Oxygen is the most abundant element on this planet. The earth's crust is 46.6% oxygen by weight, the oceans are 86% oxygen by weight, and the atmosphere is 21% oxygen by volume. Occurs as oxides, silicates, carbonates, sulfates etc., in the earth’s crust. The name oxygen comes from the Greek stems oxys, "acid," and gennan, "to form or generate." Thus, oxygen literally means "acid former." The name was introduced by Lavoisier, who noticed that compounds rich in oxygen, such as SO2 and P4O10, dissolve in water to give acids. Oxygen Two allotropes are dioxygen O2, and ozone O3 of which O2 is more stable. Isotopes 16O (99.76%), 17O (0.04%), 18O (0.20%) Fundamental building block of all vital biomolecules Gaseous O2 condenses at -183 oC to form a pale blue liquid and freezes at -219 oC to give a pale blue solid In all three phases, solid, liquid, and gas O2 is paramagnetic. Preparation of O2 Industrially Fractional distillation of liquefied air. b.p. = -183 °C (Compare with N2 -196 °C and Ar –186 °C) In the laboratory – Small amounts, rarely carried carried out o Thermal decomposition of an oxoacid salt, such as potassium chlorate, KClO3 2 KClO3 (s) heat MnO2 catalyst 2 KCl (s) + 3 O2 (g) o Electrolysis of water 2 H2O (l) 2 H2 (g) + O2 (g) o Decomposition of hydrogen peroxide 2 H2O2 (aq) catalyst 2 H2O (l) + O2 (g) Uses of O2 O2 is an inexpensive and readily available oxidizing agent. Used in many chemical industries e.g. in the making of steel (More than 66% goes towards making of steel) Acetylene – torches used in welding and cutting of metals allow for temperatures exceeding 3000 oC Used in the treatment of sewage to destroy malodorous compounds Paper bleaching to oxidize unwanted coloured compounds. Production of TiO2 – Important pigment in toothpaste and white paint Oxidation of ethene to ethylene oxide etc. Medicine Hydrogen Peroxide - The Other Hydride o Sold in drugstores as a 3% aqueous solution and as a 30% aqueous solution for industrial and laboratory uses. o Used as a mild antiseptic and as a bleach for textiles, paper, hair etc. o Pure hydrogen peroxide is a very pale blue, syrupy liquid (density = 1.45 g/cm3) and freezes at -0.4 oC and boils at 150 oC - Strong hydrogen bonds. o Pure H2O2 explodes when heated o In aqueous solutions, H2O2 behaves as a weak acid Hydrogen Peroxide - Structure Skew chain structure Weak O–O bond. Preparation of H2O2 1. large scale production: Autooxidation of an anthraquinol, e.g. 2-ethylanthraquinol 2. Lab preparation: BaO2 (s) + H2SO4 (aq) BaSO4 (s) + H2O2 (aq) Synthesis of hydrogen peroxide Redox Properties of H2O2 Ø H2O2 is both a strong oxidizing agent and a reducing agent As an oxidizing agent, oxygen is reduced from the -1 oxidation state in H2O2 to the (-2) oxidation state in H2O or (OH-). H2O2 (aq) + 2 H+ (aq) + 2 e- 2 H2O (l) Eº = + 1.78 V Oxidizing agent When H2O2 acts as a reducing agent, oxygen is oxidized from the (-1) oxidation state in H2O2 to the (0)oxidation state in O2. H2O2 (aq) O2 (g) + 2 H+ (aq) + 2 e- Eº = -0.70 V Reducing agent Hydrogen Peroxide as an oxidizing agent Hydrogen Peroxide as a Reducing agent Redox Properties of H2O2 Ø H2O2 is both a strong oxidizing agent and a reducing agent H2O2 (aq) + 2 H+ (aq) + 2 eH2O2 (aq) 2 H2O (l) Eº = + 1.78 V O2 (g) + 2 H+ (aq) + 2 e- Eº = -0.70 V H2O2 can oxidize and reduce itself and is unstable and undergoes disproportionation. 2H2O2 (l) → 2H2O (l) + O2 (g) (disproportionation) ΔH° = -196 kJ § Slow reaction but explosive reaction when catalyzed by a metal surface. § Stored in plastic bottles and a stabilizers are added. § Any substance with a standard potential in the range 0.7–1.78 V will catalyze decomposition Ozone – Trioxygen Ø Toxic, pale blue gas with a characteristic pungent odor. Ø Formation of ozone from dioxygen is extremely endothermic: 3 O2 (g) 2 O3 (g) ΔH = + 285 kJ/mol Reaction only occurs in an electrical discharge Occurs naturally in lightning Ø Ozone is a powerful oxidizing agent: O2 (g) + 4H+ (aq) + 4e- H2O (l) Eº = + 1.23 V O3 (g) + 2H+ (aq) + 2e- O2 (g) + H2O (l) Eº = + 2.07 V Ø Used in water purification, treatment of industrial waste and in organic synthesis. Structure of Ozone § Ozone is diamagnetic § O3 contains 2 equivalent O-O bonds (bond length = 128pm). § The O-O bond length is intermediate between O-O (148 pm) and O=O (121pm). § This supports the fact that O3 is resonance stabilized and the actual structure is an average of the two possibilities. N2O5 OF2 CO2 P4O10 Cl2O7 SiO2 Structure and Bonding in pblock oxides P-Block Oxides Trends in Structure and Bonding in p-Block oxides, E.g. CO2 and SiO2 In covalent compounds, oxygen generally achieves an octet configuration either by forming two single bonds or one double bond. In its oxides, CO, CO2, etc., carbon forms double bonds with oxygen, to give linear discrete covalent molecules. Unlike CO2, SiO2 is a large covalent network solid, with silica forming single bonds with oxygen. Ø Oxygen often forms a double bond to small atoms such as carbon and nitrogen because there is good p overlap between the compact p-orbitals of 2nd-row elements. Ø We can also account for these differences by evaluating the relative strengths of a double or a single bond to oxygen by carbon or silica. Trends in Structure and Bonding in p-Block oxides, E.g. CO2 and SiO2 Bond C-O C=O Si-O Si=O DH / kJmol-1 + 358 +805 +466 +638 Trends in Structure and Bonding in p-Block oxides, E.g. CO2 and SiO2 CO2 (network) CO2 (g) Need to break 4 C-O bonds and form 2 C=O bonds DHro = (4 X 358 kJmol-1) + (-2 x 805 kJmol-1) = - 178 kJmol-1 Favourable SiO2 (network) SiO2 (g) Need to break 4 Si-O bonds and form 2 Si=O bonds DHro = (4 X 466 kJmol-1) + (-2 x 638 kJmol-1) = +588 kJmol-1 Not Favourable Bond DH / kJmol-1 C—O + 358 C O +805 Si—O +466 Si O +638 Compounds of the p-block elements Bond Enthalpies – What trends do you see? Anomalous Behaviour of Period 2 Elements Summary of Main Group Elements and comparing period 2 elements with period 3 elements